Download Crystal Chemistry Atoms Electrons Quantum Mechanics Orbital

Atoms
• Element = substance that can not be chemically broken
down into smaller parts; has unique chemical properties
• Atom = smallest subdivision of an element that still retains
all of the element’s chemical properties
• Protons = positively charged particles; clustered in the
nucleus of the atom
• Neutrons = particles with a neutral charge; clustered in the
nucleus of the atom
• Electrons = negatively charged particles; located in a cloud
around the nucleus
• Atomic number = Z = number of protons in the atom; each
element has a unique number
• Isotope = same number of protons, but a different number
of neutrons
• Atomic weight = neutrons + protons + electrons
Crystal Chemistry
Chapter 3
E. Goeke, Fall 2006
• For a neutral charge, the # of
electrons = # of protons
• Ion = atom with either more or
less electrons than protons
– Anion = extra e; - charge
– Cation = missing e; +
charge
• The classic method of imaging
the relationship between
electrons & the nucleus is the
Bohr Model
– Each shell has a specific #
of electrons
– Electrons in outer shells
have more energy; the
excess energy is released as
photons when the electrons
fall into inner shells
Electrons
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
E. Goeke, Fall 2006
Quantum Mechanics
• Electrons not in a specific location, but in a zone of probability
• The electron location is described using 4 factors:
– n = principal quantum number (similar to the shell in the
Bohr model)
– l = angular momentum quantum number = designates which
kind of subshell shape; btw 0 and n-1; 1 = s, 2 = p, 3 = d, 4 =
f, and so on
– ml = magnetic quantum number = distinguishes between
orbitals of the same l value with different orientations; btw -l
and +l
– ms = spin quantum number = distinguishes between electrons
spinning to the right and those to the left; +1/2 or -1/2
• Pauli exclusion principle = no two electrons can have the same
four quantum numbers; each suborbital can have 2 electrons,
but they must spin in opposite directions
E. Goeke, Fall 2006
s orbital; increasing width &
energy with increasing n;
maximum of 2 electrons per n
p orbitals; three
orientations;
only present
when n = 2+;
max of 6e per n
d orbitals; five
orientations;
only present
when n = 3+;
max of 10e per
n
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
Mineralogy 12:041, Fall 2006, E. Goeke
E. Goeke, Fall 2006
E. Goeke, Fall 2006
Orbital Energy Levels
• Lower energy shells &
subshells are filled first
• There is a large energy
gap between the 1s and
2s orbitals, but a
smaller energy
difference between 6s
and 7s
• Note that 4s has a lower
energy level than 3d--it
will fill first
• Order shells &
subshells fill: 1s 2s 2p
3s 3p 4s 3d 4p 5s 4d 5p
6s 4f 5d 6p 7s 5f
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
E. Goeke, Fall 2006
1
s
Periodic Table
p
Ions
K
d
L
M
N
O
P
Q
f
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
Noble
gases
• Atoms tend to lose or gain electrons to fill their outer
electron shells
– Valence state = the charge on an ion; elements can have
more than 1 common valence state (e.g. Fe 2+ or Fe 3+ )
– Elements in the same column will tend to have similar
valence states
– Elements that lose electrons have a positive charge and
can be called “metals”
– Elements that gain electrons have a negative charge and
are referred to as “non-metals”
• The oxidization state of the environment plays a large role
in determining the valence state of a given ion; the more
oxygen present = more electrons removed from metals
E. Goeke, Fall 2006
E. Goeke, Fall 2006
Electronegativity
Bonding
• Electronegativity = ability of an atom within a crystal
structure to attract electrons into its outer shell
– Low EN = electron donors
– High EN = electron acceptors
– Noble gases have an EN = 0
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
• There are four kinds of bonds that we will consider in
mineralogy:
– Ionic
– Covalent
– Metallic
– Van der Waals and hydrogen bonding
• The type of bond will make a significant contribution to
the physical properties of a given mineral
• Though presented as four separate types, most bonds
actually transitional between two end members
• EN plays a role in determining what kind of bond will
form
• Charge balance must be maintained
E. Goeke, Fall 2006
E. Goeke, Fall 2006
Ionic Bonds
Covalent Bonds
• Formed from two oppositely charged ions that have either given
up or gained an electron to fill their outer shells
• Non-directional
• Dissolve easily in polar solvents (e.g. water)
• Commonly form high symmetry xtals that have moderate hardness
and density as well as high melting temperatures
• Usually poor conductors of heat & electricity
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
Mineralogy 12:041, Fall 2006, E. Goeke
E. Goeke, Fall 2006
•
•
•
•
•
•
Two atoms share 1+ electrons in order to be more stable
Very strong directional bonds
Relatively insoluble in polar solvents
High melting temperatures & high hardness
Generally form low symmetry crystals
Usually poor conductors of heat & energy
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
E. Goeke, Fall 2006
2
Covalent or Ionic?
• Elements on the right hand side of the periodic table tend to be
covalent, while elements on the left hand side are more likely to be
ionic
• Most bonds are a mix between covalent or ionic determined by the
differences in electronegativity between the two atoms
• Bonding between two atoms of the same element will produce a
pure covalent bond
• Ionic bonded crystals tend to be closely packed anions & cations
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
E. Goeke, Fall 2006
Metallic Bonds
• Hydrogen bonds = because O
has high electronegativity, it has
a greater pull on the shared e
than the H, which causes an
electrical polarity of H2O
molecules
– OH- layers result in perfect
cleavage along 1 plane
• Van der Waals bonds =
polarized atoms or molecules
result in concentrations of
positive & negative charges that
causes a slight attraction
– Weak bonds that are easy to
break
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
E. Goeke, Fall 2006
Atomic & Ionic Radii
• Size of an atom/ion will depend on the size of the nucleus
plus the number of electrons
– Those with greater numbers of electrons ~ have larger
radii
– Ions have different radii than atoms of the same
element
– As the charge becomes more positive -> fewer
electrons -> smaller radii
– Larger atomic number -> larger radii
– Radii also depend on type of bonds and how many
other ions are linked to the given ion
• Ions with similar radii are more likely to substitute for each
other
• Coordination number = CN = # of anions packed around a
cation
E. Goeke, Fall 2006
Mineralogy 12:041, Fall 2006, E. Goeke
E. Goeke, Fall 2006
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
Van der Waals & Hydrogen bonds
• Positively charged atomic nuclei share electrons amongst
themselves
• Conduct electricity & heat easily
• Low to moderate hardness
• Normally malleable & ductile
• Soluble only in acids
• Crystals usually have high symmetry
http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm
• Some elements produce hybrid s and p orbitals into sp3
orbitals during covalent bonds
– Though Si-O bonds are about 50% covalent, the bonds
produce hybrid orbitals
– The orientation of the hybrid orbitals causes the typical
silica tetrahedron arrangement
E. Goeke, Fall 2006
Ion
R(Å)
C.N. = 6
R(Å)
C.N. = 8
La+
0.74
0.92
Na+
1.02
1.18
K+
1.38
1.51
Rb+
1.52
1.61
Cs +
1.67
1.74
Fe2+
0.75
1.06
Fe3+
0.69
0.92
Nb4+
0.82
0.93
Nb5+
0.78
0.88
U6+
0.87
1.00
E. Goeke, Fall 2006
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